Lecture 12 Flashcards

1
Q

bonds

A

localised interactions that hold a structure together

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2
Q

bond direction

A

dictates the shape that things adopt

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3
Q

bond strength

A

measure of how hard it is to break a bond, energy is needed to break a bond

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4
Q

matter

A

all matter is held together by bonds, changing the bond type changes its properties

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5
Q

solid

A

strong bonds, short range, complex networks in 3D that hold things in place
lots of bonds need to be broken to remove individual components

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6
Q

ordered

A

crystalline

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7
Q

disordered

A

amorphous

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8
Q

liquid

A

less bonds, weaker
bonds are held close together but no directionality or cooperation
molecules and atoms can slide, making bonds with molecules

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9
Q

gas

A

nothing bonds to anything else, free motionp

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10
Q

phase changes

A

the conversion from one phase to another
bonds are broken because energy (heat) is added

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11
Q

ionic solids

A

rigid 3D network
moving the ions can disrupt many long range bonds
hard, high melting, brittle
insulators, don’t conduct electricity

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12
Q

covalent networks

A

rigid 3D and 2D networks
hard, high melting, brittle
insulators

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13
Q

metals

A

metallic bonding
hard (strong bonds)
ductile (atoms slide but maintain bonding)
conduct electricity

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14
Q

intramolecular

A

strong covalent bonds inside molecules

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15
Q

intermolecular

A

van der waals and hydrogen bonding between molecules

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16
Q

covalent bonding

A

doesn’t take energy to disrupt the weaker bonds
soft, low melting
requires high temperature to start decomposing (break the covalent bond)

17
Q

phase diagram

A

check notes

18
Q

ionic bonding

A

electrostatic attraction due to opposite charges between electrons
ions interacting with each other
depend on distance, not direction

19
Q

coulomb’s law

A

stronger for larger charges and smaller ions

20
Q

ionic lattice

A

non-directional, larger, 3D arrangements of the ions, ionic lattice form
the packing maximises all of the attractive (-ve to +ve) contacts, as it minimises repulsive charges
the net extra bonding outweighs the extra repulsion

21
Q

result of ionic solid

A

hard, strong, high melting solids

22
Q

ions

A

large or small
mixtures of either or both
opposite charges is the only thing that matters

23
Q

ions far apart

A

cations nucleus repels the anions nucleus and attracts its electrons
cations electrons repel the anions electrons and attract its nucleus

24
Q

medium distances

A

the gap between them is more important (the atoms are tiny, % change or r is small)
the net positive and negative charges move the ions close together

25
Q

ions close together

A

attraction doesn’t go away
the nuclei and some electrons are very close together (very strong repulsion)
the electrons that are further away from the nucleus (weaker attraction)

26
Q

ionisation

A

energy required to remove an electron from an atom

27
Q

electron gain enthalpy

A

energy required to add an electron to an ion

28
Q

electron transfer example

A

check notes

29
Q

positive energy

A

unfavourable
dangerous reaction that explodes because it gives out so much energy

30
Q

electron transfer

A

occurs because electrons ‘fall’ from a higher energy orbital to a lower energy orbital
furthered stabilised by ionic bonds, which are very strong
anion always needs a cation

31
Q

valence

A

charge an ion adopts depends on the position on the periodic table and the no of electrons used in bonding

32
Q

valence orbitals

A

outer shell is the only one involved in ionic bonding
others are too high or low in energy
these are the valence shell

33
Q

orbital energies

A

electron transfer is caused by orbital energy difference
the electron falls down that is the potential energy difference

34
Q

electronegativity

A

measure of a degree to which an atom attracts an electron in a bond
increases across a period and decreases down a group

35
Q

inverse relationship

A

as electronegativity increases, orbitals get lower in energy
number that describes the average energy of valence orbitals

36
Q

vaan aarkel triangle

A

large x = low energy, contacted orbitals
small x = radially expanded, high energy orbitals
large enough electronegativity leads to an ionic bond.